In the early 20th century, scientists were struggling to understand the structure of atoms. They had parts of the answer. The electron, which has a negative electrical charge, had been discovered earlier. They knew that the basic atom had no overall charge. Together these pieces of information made it natural to assume that the atom also contained something that carried a positive charge. Scientists guessed that since electrons are extremely small, whatever this positive something was, it must be make up most of the mass of atoms, and be much larger.

Developing a useful model of the atom

A New Zealand scientist, Ernest Rutherford, and a Danish scientist, Niels Bohr, developed a way of thinking about the structure of an atom in which an atom looks very much like our solar system. It is known as the Rutherford-Bohr Theory of Atomic Structure, and was something of a breakthrough in describing the way the atom works.

How did Rutherford and Bohr develop their theory?

Rutherford had conducted experiments in which he shot relatively large, charged particles (alpha particles) at a thin gold foil. He found that most of the particles passed directly through the foil, but some came off at odd angles, as though they had been deflected. From these results, Rutherford concluded that each atom was mostly empty space, but also contained a dense region central mass, which his alpha particles could not pass through. He also concluded that this central mass must have a positive charge, to deflect the positively charged alpha particles.

Rutherford and Bohr pictured the arrangement of the atom's parts to look like our solar system. At the center of every atom is a nucleus, which is comparable to the sun in our solar system. Electrons move around the nucleus in "orbits" similar to the way planets move around the sun.

How do the electron orbits work?

Each orbit around the nucleus represents an energy level, and electrons cannot exist in between orbits. Orbits closer to the nucleus have lower energy. If energy is added, an electron can be "excited" to jump to a higher energy level and orbit farther from the nucleus. Eventually, though, the electron will return to its original state, and the atom will give off energy equal to the difference between the two orbits.

In some materials, the energy is given off as X-rays; other materials produce specific colors of visible light, or other types of electromagnetic energy.

Each orbit can hold only a certain number of electrons. The lower-energy orbits must fill up first, if the atom is to be at its "ground" state. This is the lowest energy state and therefore most stable state.

With more research, scientists discovered that atomic structure is more complex, and that the Rutherford-Bohr model contained serious flaws.

Rutherford-Bohr Model Reexamined

Flaws in the Rutherford-Bohr Model

The Rutherford-Bohr model provided the first really useful view of the atom. It matched what scientist knew about chemical reactions and the way atoms behaved. It led to some predictions that were later proven correct. Bohr had corrected a serious flaw by recognizing that electrons had to be in orbits (energy states). But his analysis of the energy given off when an electron dropped from a higher energy orbit to a lower energy orbit didn't hold up for atoms bigger than hydrogen (the simplest atom, with only one proton and no neutrons) More work needed to be done on the model.

Improving Rutherford-Bohr - the current view of the atom

A German scientist, Erwin Schrodinger, thought the problem might be in confining the electrons to specific orbits. Other scientists had developed the idea that electromagnetic energy acted like a wave sometimes and like a particle at other times. Schrodinger thought that electrons might work the same way.

How would this idea change the Rutherford-Bohr Model?

If the electrons did behave like electromagnetic energy, we couldn't know exactly where an individual electron was. We could only know the odds (probability) of its being in a particular place.

Schrodinger replaced Bohr's well-defined orbits with probability "clouds", also known as "orbitals." He could calculate the probability that an electron would be at a particular spot in the orbital but not know for sure. In some regions of the orbital there was a high probability that an electron would be there. In others there was a low probability of electrons. The probability distributions of orbitals are sometimes shown as "lobes" extending away from the nucleus in three dimensions.

How useful was Schrodinger's idea?

Schrodinger's idea, and the equations he used to predict where electrons would be, solved problems that Bohr's couldn't. It also gave scientists a better understanding of the electron and how it behaves in chemical reactions. Schrodinger's understanding of the nature of electrons also led to research in semiconductors and other technologies on an atomic scale.

Despite its technical flaws however, the Rutherford-Bohr model is still useful because it is simple and helps people understand atomic structure.